Electrochemical Series

Electrochemical Series

The arrangement of elements in order of increasing electrode potential values is called electrochemical series. The electrochemical series is also called the activity series.

Applications of the Electrochemical Series

The main applications of the electrochemical series are:

1) Relative Strengths of Oxidising and Reducing Agents

In the electrochemical series the substances are arranged in the decreasing order of electrode potential i.e. decreasing tendency for reduction to occur or power as oxidising agent.

The elements at the top of the table have maximum tendency to get reduced and consequently they will act as good oxidising agents.

The substances at the bottom of the table have lower electrode potential values, therefore they have least tendency to get reduced. Consequently they may be oxidised and act as good reducing agents.

 

Standard electrode potential of fluorine is the highest in the table indicating that the fluorine gas (F₂) has maximum tendency to get reduced to fluoride ions (F¯) and therefore, fluorine gas is the strongest oxidising agent and fluoride ion is the weakest reducing agent.

Lithium has the lowest electrode potential in the table indicating that lithium has the maximum tendency to get oxidised and hence lithium metal is the most powerful reducing agent or lithium ion is the weakest oxidising agent in aqueous solution.

The substances which have lower electrode potentials are stronger reducing agents while those which have higher electrode potentials are stronger oxidising agents

2) Calculation of the E.M.F. of the cell

The E.M.E. of the cell which is the difference between the reduction potential of the cathode and anode is determined by the following steps:

Step 1: Write the two half cell reactions in such a way that the reaction taking place at the left hand electrode is written as an oxidation reaction and that taking place at the right hand electrode is written as reduction.

Step 2: Multiply by suitable number to equate the number of electrons in the two equations. However, it may be noted that electrode potential values, E^ø are not multiplied.

Step 3: The e.m.f. of the cell (E_cell) is equal to the difference between the standard electrode potential of the cathode and the standard electrode potential of the anode. Thus,

E^ø= E^ø(cathode) – E^ø(anode)

The electrode potentials of both the electrodes are taken to be reduction potentials.

Step 4 :If the e.m.f. of the cell is +ve, the reaction is feasible in the given direction. But if e.m.f. of the cell is -ve, the cell reaction is not feasible in the given direction. The reaction must be occurring in the reverse direction. Thus, to get positive value for the e.m.f. of the cell, the electrodes are reversed.

 

3) Predicting feasibility of the reaction

A redox reaction is feasible only if the species which has higher reduction potential is reduced i.e. accepts the electrons and the species which has lower electrode potential is oxidised i.e., loses the electrons. Otherwise, a redox reaction is not feasible.

The species to release electrons must have lower electrode potential as compared to the species which is to accept electrons.

The electrochemical series gives the increasing order of electrode potential (reduction) of different electrodes on moving down the table. The species to accept the electrons (getting reduced) must be lower in the electrochemical series as compared to the other which is to lose electrons (getting oxidised).

Example 

Cu²⁺ (aq) +2Ag(s) ——-> Cu(s) +2Ag⁺ (aq)

E° value of Cu +0.34 V and the E value of Ag + 0.80 V. Since the electrode potential of Ag is more than that of Cu, this means that silver has greater tendency to get reduced in comparison to copper. Thus, the reaction

Ag⁺ (aq) +e¯ ———> Ag(s)

occurs more readily than the reaction

Cu²⁺ (aq) + 2e¯ ———->Cu(s)

Similarly, since reduction potential of Cu is less than that of Ag, this means that Cu will be readily oxidised in comparison to Ag. Thus, the reaction

Cu(s) ——->Cu²⁺(aq) + 2e¯

 

occurs more readily than the reaction

Ag(s) ——–> Ag⁺ (aq) +e¯

Therefore, silver will be reduced and copper will be oxidised and the above reaction is not feasible.

Cu(s) +2Ag⁺(aq)———>Cu²⁺(aq)+ 2Ag(s) can occur

Similarly, we can predict that iron will reduce Cu²⁺ to Cu, zinc will reduce Pb²⁺ to Pb, zinc will also reduce Cu²⁺ to Cu, etc and the following redox reactions can occur:

Fe +  Cu²⁺ ——-> Fe²⁺ + Cu

Zn + Pb²⁺ ——–>Zn²⁺ + Pb 

4) To predict whether a metal can liberate hydrogen from acid or not

Metals like zinc, magnesium and nickel can liberate hydrogen from the acids like HCl, H₂SO₄, etc. while metals like copper and silver cannot do so.

Only those metals can liberate hydrogen from the acid which have negative values of reduction potentials i.e. -E° value.

Hydrogen will have greater tendency to get reduce (accept electrons) and the metal can lose electrons (get oxidised) and hydrogen gas is liberated. These metals are also called active metals.

M —–> M⁺(aq) + e¯

(Monovalent)

H+(aq) +e‾ ———> H

Metals like copper and silver have +E° values i.e. electron accepting tendencies. Their atoms are not in a position to lose electrons to H⁺ ions of the acid. Therefore, hydrogen gas is not liberated.

If the standard electrode potential of an electrode is greater than zero (i.e.+ve), then its reduced form is more stable as compared to hydrogen gas.

If the standard electrode potential is negative, then hydrogen gas is more stable than the reduced form of the species.